2. ATOMIC THEORY Building blocks of matter

3. Who are these men? In this lesson, we’ll learn about the men whose quests for knowledge about the fundamental nature of the universe helped define our views.

4. DEMOCRITUS IN 400 BC, DEMOCRITUS SAID:  All matter is made of tiny particles called “atomos”  Disputed by Aristotle

5. Why? Eminent philosophers of the time, Aristotle and Plato, had a different idea. They favored the earth, fire, air and water theory of matter. They were more popular, so the atomos idea was buried for approximately 2000 years.

6. Atomos   To Democritus, atoms were small, hard particles like marbles with different shapes and sizes.   Atoms were infinite in number, always moving and capable of joining together.

7. For the next 2000 years…  Alchemists tried to make gold from other metals. UNTIL… UNTIL… 1808 JOHN DALTON’S NEW ATOMIC THEORY

8. LAWS FROM ATOMIC THEORY (from last chapter)  1. Law of conservation of matter—the mass of the reactants before the reaction equals the mass of the products after.  2. Law of definite proportions—Every sample of the same compound has the same mass ratio of component elements.  3. Law of multiple proportions—In a series of compounds of the same two elements, the ratio of an element in one compound to another is also a small, whole number.

9. Dalton’s Model  In the early 1800s, English chemist John Dalton performed careful experiments that eventually led to the acceptance of the idea of atoms.

10. Dalton’s atomic theory  1. All matter is made up of atoms  2. Atoms of the same element are alike.  3. Atoms of different elements are different.  4. Compounds have a definite composition by weight and combine in small whole number ratios.  5. Atoms cannot be subdivided.

11. .  This theory became one of the foundations of modern chemistry.

12. Thomson’s Plum Pudding Model  In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles.

13. Thomson Model  His model of the atom is sometimes called the “Plum Pudding” model.  Atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding.

14. Thomson Model  Thomson studied passing an electric current through a gas.  As the current passed through, it gave off rays of negatively charged particles.

15.  This surprised Thomson, because the atoms of the gas were uncharged. Where had the negative charges come from? Where did they come from?

16. He concluded that the negative charges came from within the atom. A particle smaller than an atom had to exist. was The atom was divisible! Thomson called the negatively charged “corpuscles,” today known as electrons. Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom. But he could never find them.

17. Ernest Rutherford  In 1908, the English physicist Ernest Rutherford was hard at work on an experiment that had little to do with unraveling the mysteries of the atomic structure.

18.  Rutherford’s experiment involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)

19. Rutherford –Most of the positive particles passed through the gold atoms in the foil without changing course at all. –Some of the positive charges did bounce away from the gold sheet as if they had hit something solid. He knew that like charges repel.

21. Rutherford  The gold atoms in the sheet were mostly empty space. Atoms were not a plum pudding.  Atom has a small, dense, positively charged center that repelled the positive “bullets.”  He called the center of the atom the “nucleus”  The nucleus is tiny compared to the atom as a whole.

22. http://chemmovies.unl.edu/ChemAnime/RUTH ERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/RUTH ERFD/RUTHERFD.html

23. Atomic Particles  Electron—discovered by Thomson in 1890’s –Robert Millikan—determined the charge of an electron in 1909 w/ oil drops  Proton—discovered by Rutherford in 1911  Neutron—discovered by James Chadwick in 1932

24. Particles and Charge Mass Charge Location Mass Charge Location  Proton (p + )  Neutron (n o )  Electron (e - ) 1 amu +1 nucleus 1 amu +1 nucleus 1/1840 amu -1 electron cloud 1/1840 amu -1 electron cloud 1 amu 0 nucleus 1 amu 0 nucleus

25. How Atoms Differ Dalton said that all atoms of an element are alike, but we know that is not completely true. So what is alike? 1.All atoms of the same element have the same number of protons. 2.If the atom is neutral, that means they also have the same number of electrons. 3.The number of neutrons, however, can vary.

26. How Atoms Differ Atoms with the same atomic number are the same element, but they may have different numbers of neutrons.  Atoms of the same element with a different number of neutrons are called isotopes.  Atoms of the same element with a different number of electrons than protons are called ions.

27. The Chemists’ Shorthand: Atomic Symbols for Isotopes K  Element Symbol 39 19 Mass number  Atomic number  #p + + #n o #p + or #e - Mass # - Atomic # = #n o

28. Atomic Masses  Elements occur in nature as mixtures of isotopes  Atomic mass is the weighted average of all isotopes for an element.  Carbon =98.89% 12 C 1.11% 13 C <0.01% 14 C  Carbon atomic mass = 12.01 amu

29. MASS NUMBER AND AVERAGE ATOMIC MASS Atomic masses are based on CARBON. The atomic mass unit is 1/12 of the mass of one carbon atom. How do we calculate average atomic mass? Multiply the % times the mass for each isotope, then add them together.

30. Average atomic mass  Calculate the average mass of isotopes of neptunium with: 50.0% at 238.05 amu 29.4% at 235.1 amu 20.6% at 237.98 amu (.500 x 238.05) + (.294 x 235.1) + (.206 x 237.98) = 237.17amu

31. Another problem: Calculate the average atomic mass of calcium with these isotopes: 28.4% at 40.06 amu 34.1% at 41.02 amu 22.8% at 40.89 amu 14.7% at 39.98 amu (.284x40.06)+(.341x41.02)+(.228x40.89)+(.147x39.98)40.56

32. One more for Arsenic 35.1% of 74.9 amu 18.6% of 74.2 amu 46.3% of 75.02 amu 74.83 amu

33. Atomic Mass  Atomic mass is the weighted average of all of the known isotopes of an element, so will always be shown as a decimal number.

34. Covalent Chemical Bonding  The forces that hold atoms together in compounds. Covalent bonds result from atoms sharing electrons between nonmetal atoms.  Molecule: a collection of covalently-bonded atoms.  Atom: representative particle for a monatomic element

35. Ionic Chemical Bonding Cation: A positive ion Mg 2+, NH 4 + Anion: A negative ion Cl , SO 4 2  Ionic Bonding: Force of attraction between oppositely charged ions. Smallest particle called a formula unit.

36. Stupendous Seven

37. Periodic Table Elements classified by: properties properties atomic number atomic number Groups (vertical columns)—also called families 1A = alkali metals 1A = alkali metals 2A = alkaline earth metals 2A = alkaline earth metals 7A = halogens 7A = halogens 8A = noble gases 8A = noble gases Periods (horizontal rows)

38. Periodic Table Antoine Lavoisier, 1790’s made first list of known elements, 23 total. By 1870, there were 70! John Newlands, 1864—Law of Octaves: When element were placed in order of increasing atomic mass, every 8 th element repeated properties. Lothar Meyer, 1869—Periodic table based on physical characteristics only and increasing atomic mass. Dmitri Mendeleev, 1869—Periodic table based on physical and chemical characteristics and increasing atomic mass. Predicted new elements. Henry Moseley, 1913—Modern periodic law based on subatomic particles: There is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number (protons).

39. Periodic Trends

40. Groups Periods Groups Periods

41. Periodic Trends